Cold Pack Chemistry: Exploring Endothermic and Exothermic Reactions
OverviewHow do "instant" cold packs get cold when they are stored at room temperature, unlike a regular ice pack which must be stored in the freezer? In this lesson plan, students will explore several endothermic and exothermic reactions, and use their observations to choose the chemical reaction that best fits the design constraints for their own chemical cold pack.
- Define the criteria and constraints of an engineering problem (designing an instant cold pack)
- Identify the chemical reaction that best meets the criteria and constraints
NGSS AlignmentThis lesson helps students prepare for these Next Generation Science Standards Performance Expectations:
- MS-PS1-6. Undertake a design project to construct, test, and modify a device that either releases or absorbs thermal energy by chemical processes.
- MS-ETS1-1. Define the criteria and constraints of a design problem with sufficient precision to ensure a successful solution, taking into account relevant scientific principles and potential impacts on people and the natural environment that may limit possible solutions.
Science & Engineering Practices
Analyzing and Interpreting Data. Analyze data to define an optimal operational range for a proposed object, tool, process, or system that best meets criteria for success.
Disciplinary Core Ideas
PS1.A: Structure and Properties of Matter. Some chemical reactions release energy, others store energy.
ETS1.A: Defining and Delimiting Engineering Problems. The more precisely a design task's criteria and constraints can be defined, the more likely it is that the designed solution will be successful. Specification of constraints includes consideration of scientific principles and other relevant knowledge that are likely to limit possible solutions.
Energy and Matter. The transfer of energy can be tracked as energy flows through a designed or natural system.
For the entire class:
- Optional: at least one instant cold pack for a demonstration
- Access to a sink
- Paper towels
- Well-ventilated area or fume hood
Each group will need:
- 50 mL (or larger*) beaker
- Cup or other container for pouring water/vinegar
- Spatula or spoon for chemicals
- Access to a scale
- Partial immersion thermometer
- Optional: stirring rod
- Disposable gloves
- Chemical safety goggles
- Ammonium chloride
- Magnesium sulfate
- Sodium bicarbonate (baking soda)
- Tap water
* A 50 mL beaker will allow you to immerse the thermometer bulb in about 20 mL of water. If you use larger beakers, you will need to make sure you use enough water to immerse the thermometer bulb, and increase the amounts of the other reactants proportionally so you maintain the same concentrations.
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Background Information for TeachersThis section contains a quick review for teachers of the science and concepts covered in this lesson.
Instant cold packs (Figure 1) are "ice packs" that, unlike regular ice or gel packs, do not need to be stored in a freezer. Inside the outer bag, water and another chemical (typically ammonium nitrate) are separated by a second bag or tube. When you squeeze the outer bag, the inner bag breaks, allowing the ammonium nitrate to mix with the water. This results in an endothermic chemical reaction—one which absorbs heat, causing a local decrease in temperature. This is in contrast to an exothermic chemical reaction, which releases heat, causing a local increase in temperature.
an instant cold pack
Figure 1. Two instant cold packs.
Designing an instant ice pack is both a chemistry problem and an engineering problem. There are criteria and constraints around the design of the ice pack. For example, how big should it be? How cold should it get? How cold is too cold? How long will it stay cold? In turn, these questions can be used to inform decisions about which chemicals to use and the corresponding chemical reaction. How should the reaction be activated? What is the proper ratio of reactants? Are the chemicals safe if the bag breaks and they spill? In this lesson plan, your students will explore different chemical reactions as they pose questions like these themselves (and seek answers). We recommend the chemical reactions shown in Table 1, although you could substitute or add others. Results from Science Buddies staff testing of these reactions are show in Figure 2.
|Chemicals||Chemical Equation||Reaction type||Safety Notes|
|Calcium chloride and water||Exothermic||Can get very hot (about 200°F/93°C). Do not touch with bare hands!|
|Ammonium chloride and water||Endothermic|
|Magnesium sulfate and water||Endothermic|
|Sodium hydrogen carbonate (baking soda) and vinegar||Endothermic||May overflow and spill if beaker is too small.|
This graph shows a plot of minimum or maximum temperature in degrees F versus concentration in g/mL for four different chemical reactions: ammonium chloride and water, calcium chloride and water, magnesium sulfate and water, and baking soda and vinegar. The ammonium chloride graph is concave up with a minimum of 36 degrees at a concentration of 0.3 g/mL. The calcium chloride graph is concave down with a maximum of 182 degrees at a concentration of 0.9 g/mL, however the right-hand side of the curve is truncated because higher concentrations were not tested. The magnesium sulfate graph is concave up with a minimum of 58 degrees at a concentration of 0.7 g/mL. The baking soda graph is very slightly concave up with a minimum of 60 degrees at a concentration of 0.3 g/mL.
Figure 2. Experimental data for maximum/minimum temperature vs. concentration for the chemical reactions shown in Table 1.